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      "body": "<html>\n<h2>Intro</h2>\n<p>In my previous <a href=\"https://steemit.com/corrosion/@rahim.rahman/how-corrosion-occurs-basic-electrochemistry\">post</a>, I have often used 2 dissimilar metals as example when describing about corrosion cell (electrochemical cells), where a complete circuit is formed when (i) : there is electrical contact between anode and cathode via a metallic path (wire) and (ii): ions can flow through the electrolyte. However, corrosion cell can also occur &nbsp;in a single (isolated) metal. For example, household gates or fences are often built with a single metal material (mild steel) and overtime, they will corrode although without coupled with another metal.&nbsp;</p>\n<h2>Corrosion Cell - Single Metal</h2>\n<p>Let us begin with comparison of diagram between the 2 corrosion cells. &nbsp;They are shown below:</p>\n<p><img src=\"https://i.imgsafe.org/2b/2b20e6c0e0.jpeg\" width=\"490\" height=\"270\"/></p>\n<p>A= anode, &nbsp;&nbsp;C=cathode, &nbsp;&nbsp;E=electrolyte, &nbsp;&nbsp;e=electron, &nbsp;&nbsp;MP=metallic path</p>\n<p>As can be seen, they are not that different. Corrosion cell occurring within a single metal can arise due to many factors. Do take note that, in reality there is no single material in this world that is perfectly uniform in characteristics. The major factors that contribute to this phenomenon can include:</p>\n<ol>\n  <li>Differences in microstructure of the material &nbsp;(i.e- differences in grain boundary orientation, many phases within the material)</li>\n  <li>Foreign material within the material (i.e - inclusions like oxides, sulfides which promotes anodic/cathodic site)</li>\n  <li>Differential aeration (i.e- lower oxygen concentration promotes anodic site [such as metal traps within soil])</li>\n  <li>Heat effect (i.e - welding job, metal heat treatments)</li>\n  <li>Mechanical work (i.e-strained areas tend to become anodic site)</li>\n</ol>\n<p>Let's take for example, a piece of iron that is left outside and exposed to moisture. Due to exposure with moisture (i.e-droplet of water), the iron surface that is in contact with water droplet becomes anode while the iron at the &nbsp;edge of the water droplet becomes cathode. Electron will move from the anodic site to the cathodic site within the metal itself. The reactions for this phenomenon &nbsp;are:&nbsp;</p>\n<ol>\n  <li>Oxidation: &nbsp;&nbsp;Fe --&gt; Fe(2+) + 2e(-) &nbsp;</li>\n  <li>Reduction: &nbsp;O(2) + 2H(2)O + 4e(-) &nbsp;--&gt; 4OH(-) &nbsp;[Just outside water droplet]</li>\n</ol>\n<p>The hydroxide ions [4OH(-)] can move towards the water droplet and react with Fe(2+) ions to form iron (II) hydroxide:</p>\n<ol>\n  <li>Iron (II) hydroxide: &nbsp;Fe(2+) + 2OH(-) --&gt; Fe[(OH)(2)]</li>\n</ol>\n<p>The iron hydroxide is then oxidized to form rust. For ease of visualisation, we may refer to the below figure:</p>\n<p><img src=\"https://i.imgsafe.org/2b/2bb7943ada.jpeg\" width=\"587\" height=\"336\"/></p>\n<p>I &nbsp;hope this post will allow others to understand more about corrosion phenomenon in general.</p>\n<p>Stay tuned for more posts.</p>\n<p>Cheers,</p>\n<p>@<a href=\"https://steemit.com/@rahim.rahman\">rahim.rahman</a></p>\n<p><br></p>\n</html>",
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      "body": "<html>\n<h2>Intro</h2>\n<p>Based on <a href=\"https://steemit.com/corrosion/@rahim.rahman/factors-influencing-corrosion\">EMF series</a> and <a href=\"https://steemit.com/corrosion/@rahim.rahman/how-corrosion-occurs-basic-electrochemistry\">galvanic series</a>, we can notice that the common parameter for both series is potential (voltage). One may wonder, how about current? Does current play a role too in corrosion? As a matter of fact, it does but current alone is not the main parameter. The main parameter is called <strong>current density</strong>, which is a measure of <code><em>current per unit area</em></code>.&nbsp;</p>\n<h2>Current Per unit Area</h2>\n<p>Current per unit area plays a major role in affecting corrosion rate. As can be seen from my previous <a href=\"https://steemit.com/corrosion/@rahim.rahman/how-corrosion-occurs-basic-electrochemistry\">post, </a>current flows from cathode to anode, which is opposite to the flow of electron. Let's say we have a cathode which is bigger than the anode. As the current flows from big area of cathode to the small area of anode, the total current has to be distributed to the surface area belonging to the small anode. Since the large amount of &nbsp;current has to flow to the anode, the current per unit area is high at anode. This will force the anode to accept high flows of current. The effect is concentrated over a small area, which leads to higher corrosion rate.&nbsp;</p>\n<h2>Evans Diagram</h2>\n<p>Remember that in <a href=\"https://steemit.com/corrosion/@rahim.rahman/corrosion-pourbaix-diagram\">Pourbaix</a> diagram, potential is plotted against pH. For Evans diagram, the y-axis remains as potential but the x-axis is now the current density in logarithmic scale. Typical Evans diagram for zinc in acid solution is shown as below:</p>\n<p><img src=\"https://i.imgsafe.org/fe/fe91499468.jpeg\" width=\"534\" height=\"375\"/></p>\n<p>From the above diagram, notice that there are 4 slopes existed, 2 comes from Zn &amp; 2 comes from H(2). The 2 slopes from each element corresponds to oxidation (sloping up) reaction and reduction (sloping down) reaction. Once the Zn is connected/immersed in hydrogen, both of them experience <strong>activation polarization</strong> along their lines. When the 2 lines meet, the <strong>corrosion potential (E_corr)</strong> &amp; <strong>corrosion current density &nbsp;(i_corr)</strong> will be known. From here, the corrosion rate of zinc can be estimated using current density value (i_corr). Do take note that the slope of Evans diagram is usually obtained experimentally, although Tafel equation may used to &nbsp;construct the slopes (not explained due to complexities). For everybody's info, the experimentation work to obtain Evans diagram involves conducting potentiodynamic polarization measurement, using &nbsp;a machine called potentiostat. Please refer <a href=\"ftp://185.72.26.245/Astm/1/Section%2003/ASTM0302/PDF/G59.pdf\">here</a> for further reading on the details. &nbsp;</p>\n<p>Apart from referring to Evans diagram, E_corr &amp; i_corr can also be obtained from equations. For example, the equations for zinc and hydrogen will be in the form as followed:</p>\n<p><img src=\"https://i.imgsafe.org/ff/ff6d46a63c.jpeg\" width=\"213\" height=\"59\"/>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;<img src=\"https://i.imgsafe.org/ff/ff6d589e58.jpeg\" width=\"224\" height=\"52\"/></p>\n<p>Where B and i(o) are constants for each cell. By equating V(H) &amp; V(Zn), i_corr can then be obtained.</p>\n<h2>Polarization (change in voltage)</h2>\n<p>Now that we are familiar with Evans diagram, we know that corrosion rate can be reduced when i_corr is lowered. This can be achieved by 2 methods, either by increasing the slope of anodic reaction (anodic inhibitor) or increasing the slope in -ve direction for cathodic reaction (cathodic inhibitor).&nbsp;</p>\n<h3>Anodic Inhibitor</h3>\n<p>We can use anodic inhibitor to increase the slope of anodic reaction. Some examples of anodic inhibitor would be:</p>\n<ol>\n  <li>&nbsp;Oxidizing anodic inhibitor: Chromate, Nitrite (can passivate steel without oxygen)</li>\n  <li>&nbsp;Non-Oxidizing anodic inhibitor: Phosphate, Tungstate &amp; Molybdate (required oxygen to passivate steel)</li>\n</ol>\n<p>The effect of changing the anodic reaction can be seen as below:</p>\n<p><img src=\"https://i.imgsafe.org/00/00139a55f9.jpeg\" width=\"611\" height=\"326\"/></p>\n<h3>Cathodic Inhibitor</h3>\n<p>Instead of changing the slope of anodically, we can also change the slope cathodically to reduce the corrosion current density as shown in the figure below:</p>\n<p><img src=\"https://i.imgsafe.org/00/000d1a6498.jpeg\" width=\"620\" height=\"315\"/></p>\n<p>Examples of cathodic inhibitor are arsenic, bismuth and antimony compounds. However, there are drawbacks with this method. The compounds can cause hydrogen blistering of steel and also easier to experience hydrogen embrittlement. This may occur because cathodic inhibitor prevents hydrogen atom formation, which in turn causing concentration of hydrogen atoms at the surface of the metals which will then absorbed into the steel. Hydrogen atoms then combined to form hydrogen molecules inside the steel. Since hydrogen molecules do not &nbsp;diffuse through steel, they tend to gather at voids or defects and the creating pressure which may reach pressure up to 7k &nbsp;MPa to cause pitting/damage.</p>\n<h3>Anodic Protection</h3>\n<p>Another type of corrosion control is by maintaining a metal in the passive region by means of externally applied anodic current. By using an external power supply, it is possible to achieve passivity by altering the potential. The Evans diagram for such metals will typically has a S-shape like this:</p>\n<p><img src=\"https://i.imgsafe.org/15/1557e8aae4.jpeg\" width=\"557\" height=\"383\"/></p>\n<p>Typical example for this would be a steel storage tank for sulphuric acid (example in a chemical plant).</p>\n<h3>Point to take note</h3>\n<p>As of now, we have covered on brief explanation on Evans diagram &amp; Pourbaix diagram. One important point to take note is that, Pourbaix diagram won't able to tell the rate of corrosion. It gives us information on whether a reaction is feasible (possibility of corrosion). &nbsp;Meanwhile, Evans diagram is able to tell us the corrosion rate. Do take note that Evans diagram deals with the kinetic aspect of corrosion &nbsp;while Pourbaix diagram deals with thermodynamic aspect of corrosion.</p>\n</html>",
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      "body": "<html>\n<h2>Intro</h2>\n<p>In previous <a href=\"https://steemit.com/corrosion/@rahim.rahman/factors-influencing-corrosion\">post</a>, it has been mentioned that among the factors that can affect corrosion are pH and potential (voltage). Under certain pH and potentials, a metal can become unstable. Knowing these 2 parameters will be beneficial for anyone wishing to estimate whether corrosion can occur or not under different pH and potentials.&nbsp;</p>\n<h2>Pourbaix Diagrams</h2>\n<p>A useful diagram for corrosion estimator based on pH and potential is called Pourbaix diagram. Pourbaix diagram was created by a Belgium corrosion scientist named Pourbaix. It is basically a potenial-pH diagram, which <strong>represents the stability of a metal as a function of potential and pH</strong>. The diagram is constructed based on Nernst equations and &nbsp;metals &amp; its species solubility in equilibrium state. In this diagram, we can see 3 different regions which are:</p>\n<ol>\n  <li>Immune region (thermodynamically stable, no corrosion)</li>\n  <li>Corrosive region (active state)</li>\n  <li>Passive region (forming passive layers, inhibiting corrosion)</li>\n</ol>\n<p>Before we get into the details of the diagram, let us find out the diagram's objectives:</p>\n<ol>\n  <li>To know which pH and potential changes will prevent corrosion</li>\n  <li>To estimate corrosion product compositions at different pH &amp; potential</li>\n  <li>To know the directions of various reactions at different pH &amp; potential</li>\n</ol>\n<p>Ok, now let's get into the details/characteristics of Pourbaix diagram:</p>\n<ol>\n  <li>Hydrogen and oxygen lines are drawn in dotted line.</li>\n  <li>The Pourbaix diagram is shown for equilibrium conditions at 25 deg C.</li>\n  <li>The concentration of metal ions is assumed to be 10^-6 mol. (Note: lower concentration shouldn't cause corrosion)</li>\n  <li>X-axis is showing pH value while Y-axis is showing redox potential vs SHE value. (SHE: standard hydrogen potential)</li>\n  <li>Horizontal lines represent electron transfer reaction which are pH independent.</li>\n  <li>Vertical lines doesn't represent electron transfer reaction, which are pH dependent.&nbsp;</li>\n  <li>Sloping lines give redox potentials of a solution in equilibrium with hydrogen and oxygen, respectively.</li>\n  <li>Oxidizing power increases with increasing potential.</li>\n</ol>\n<h3>Pourbaix for Water</h3>\n<p>The redox reactions for water can be referred from EMF series. One of the reactions is (1/2)O(2) +2H(+) + 2e(-) --&gt; H(2)O, which&nbsp;has a potential of 1.23V. By using this potential value, knowledge from pH applicability of Nernst equation, &nbsp;the Nerst equation becomes: E = 1.23 - 0.059pH.</p>\n<p>Another redox reaction for water is 2H(+) +2(e) --&gt; H(2) , which&nbsp;has a potential of 0V . Using similar technique as above, the Nerst equation becomes: E = &nbsp;- 0.059pH.</p>\n<p>These 2 equations are then plotted on the Pourbaix diagram. The diagram is shown as below:</p>\n<p><img src=\"https://i.imgsafe.org/e6/e683601604.jpeg\" width=\"422\" height=\"352\"/></p>\n<p>The info for the above diagram is such that:</p>\n<ul>\n  <li>above line (a), water is stable and any H(2) present is oxidised to water</li>\n  <li>below line (b), water is stable and any O(2) present is reduced to water</li>\n  <li>above line (b), water is unstable and oxidize to give O(2)</li>\n  <li>below line (a), water is unstable and reduce to H(2)</li>\n</ul>\n<p>Now that we have formed the lines for water, there is one thing left to do, which is to add line for metal. For iron, the Pourbaix diagram looks like this:</p>\n<p><img src=\"https://i.imgsafe.org/e7/e793279af8.jpeg\" width=\"421\" height=\"431\"/></p>\n<p><br></p>\n<p>The lines are formed by using the same method as mentioned above in addition with Gibbs free energy calculation(for simplicity, it is not shown here). Region of stability for iron is at the bottom part of the diagram, which lies about below -0.6V (some other sources mentioned as below -0.44V). Several important observations from the diagram are:</p>\n<ol>\n  <li>Iron will be unstable in water, no matter what are the pH and potential. That is, iron will corrode in water.&nbsp;</li>\n  <li>Iron can be protected by passivation [coating with Fe(2)O(3)] or cathodic protection [coupled with more active metal that has larger negative potential].</li>\n</ol>\n<p>We can relate real world phenomenon with Pourbaix diagram. For example, bridge's components that are submerged in water or mud will tend to corrode terribly than other bridge components in air. This is because the mud or water has potential(voltage) that is close to H(2) line, where iron,Fe can corrode to Fe(2+). The corroded and non-corroded regions for iron,Fe in Pourbaix diagram is shown below:</p>\n<p><img src=\"https://i.imgsafe.org/e8/e865040e4a.jpeg\" width=\"509\" height=\"308\"/></p>\n<p>OR in simple form:</p>\n<p><img src=\"https://i.imgsafe.org/e8/e860a3bd7d.jpeg\" width=\"526\" height=\"345\"/></p>\n<p><br></p>\n<p>The Pourbaix diagram do have some limitations such as follows:</p>\n<ol>\n  <li>The diagram is only applicable for 25 deg C. Can't predict corrosion at higher temperature.</li>\n  <li>It tells whether corrosion happens or not but not the corrosion rate.</li>\n  <li>Only applicable for pure metals. Not applicable for alloys or impurities effect. Hence, not really practical for engineers. &nbsp;</li>\n</ol>\n<p>&nbsp;That's all about Pourbaix diagram. I hope readers may learn something from this post.&nbsp;</p>\n<p>Thank you.</p>\n<p>Cheers,</p>\n<p>@<a href=\"https://steemit.com/@rahim.rahman\">rahim.rahman</a></p>\n</html>",
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      "body": "@@ -3738,16 +3738,528 @@\n %22/%3E%3C/p%3E%0A\n+%3Cp%3E1 important observation is that, we can use EMF series to predict the potentials (voltage) for the image above. The voltage from the image showed a reading of 1.067 V. Let's compared this to values  in EMF  series. According to EMF series, zinc shows a value of -0.763 V while copper shows a value of  +0.34 V. By subtracting 0.34 & (-0.763), this will yield 1.103 V.  As you can see, there isn't much difference in value between the experimentation work and the EMF series. %3C/p%3E%0A\n %3Ch1%3ENern\n",
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      "body": "@@ -2796,34 +2796,137 @@\n but \n-at a quicker rate than the\n+with a higher driving force for the electrochemical reactions.  The corrosion effect will be greater (looking more bad) than\n  cas\n",
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      "body": "<html>\n<h2>Intro</h2>\n<p>Corrosion is quite a complicated phenomenon. Lots of expertise from various fields (engineering, science, physics) have been working together to understand more on it and trying to prevent as well as control the corrosion from happening. It's among the active research area to be conducted by universities and industries. For those who are interested in knowing about the basics of corrosion, I will try to share some fundamentals on corrosion in this post. Hope we all can &nbsp;learn something and appreciate the science behind corrosion.</p>\n<h2>Potential Difference</h2>\n<p>When corrosion occurs, there will be chemical reactions (i.e-oxidation,reduction). Often the case, chemical reactions produce heat. However, the heat is quite difficult to detect since the rate of the chemical reactions are quite slow. Instead of measuring the heat, it is more practical to deal with electrical energy that involves with corrosion. In other words, energy that flows during corrosion reactions is viewed in the form of electrical energy.&nbsp;</p>\n<p>By using voltmeter, we can measure the electrical potential (voltage) between 2 points, which is the energy differences. By using this method, we can find the potential difference between an anode and a cathode (i.e any 2 different metals), where the anode is in higher energy state than the cathode. In science, energy always travel from high energy state to lower energy state which means, during corrosion; the electrons will flow from high energy state (anode) to low energy state (cathode).</p>\n<p>A voltmeter typically has 2 wires of different colours, a red and a black wires. The convention is that, the black wire should be connected to the negative terminal while the red wire connects to the positive terminal. By doing this, a positive value will be displayed by the voltmeter. Now, imagine if we have a zinc plate and a copper plate connected by electrochemical cell. If the black wire is connected to the zinc while the red wire is connected to the copper and the voltmeter is displaying a positive value, this means the zinc is the anode while the copper is the cathode. For visualisation purposes, here is the diagram.</p>\n<p><img src=\"https://i.imgsafe.org/b8/b8e9494aa1.jpeg\" width=\"539\" height=\"328\"/></p>\n<p>As you can see, the voltmeter is displaying a positive value, meaning that the black wire is connected to the negative terminal (anode) while the red wire is connected to the positive terminal (cathode).&nbsp;</p>\n<p>P/S: Note that, the electrolytes usually will be in in 1 mol ions concentration.</p>\n<h2>EMF Series</h2>\n<p>When dealing with corrosion topic, EMF series is one of the fundamentals. The EMF series or also called as electromotive force series is quite similar to <a href=\"https://steemit.com/corrosion/@rahim.rahman/how-corrosion-occurs-basic-electrochemistry\">galvanic series</a>. It lists metals according to their oxidation potentials (active metals having larger negative potentials while passive metals not so negative). The differences are that:</p>\n<ol>\n  <li>The list is only for pure metals. For example, it doesn't has stainless steel in the list because stainless steel is not in pure form. It consists of several elements. (i.e - iron (Fe), chromium (Cr), nickel(Ni) and etc..)</li>\n  <li>EMF series has only 1 list. The list is generated based on experiment conducted in solutions of standards 1 mol ion concentration. EMF series, only exist for a single environment while for galvanic series, it has many lists, depending on &nbsp;the environment of the metals.</li>\n</ol>\n<p>The &nbsp;EMF &nbsp;series is shown here as below:</p>\n<p><img src=\"https://i.imgsafe.org/b9/b98ecbeb1a.jpeg\" width=\"577\" height=\"397\"/></p>\n<h1>Nernst Equation</h1>\n<p>As have been mentioned before, EMF series is only applicable to pure metals in 1 mol solutions. If the concentrations is different than 1 mol, Nerst equation has to be used to obtain the right potentials. At 25 deg C (room temperature). Nerst equation is in the form as below:</p>\n<p>&nbsp;<img src=\"https://i.imgsafe.org/b9/b9d0e6d703.jpeg\" width=\"298\" height=\"75\"/></p>\n<p>Where V2 and V1 are the electrodes potentials, n is the number of electrons in half cell reactions, M1 and M2 are the concentrations of ions in mol.</p>\n<h2>pH</h2>\n<p>Another factor that can affects corrosion is pH. As many people may have already known, pH is a measure of the acidity or alkalinity of a solution. In fact, pH is a short form of <em>potential hydrogen. </em>In actuality, pH refers to the hydrogen ion activity within a solution. Higher hydrogen ion activity corresponds to lower pH &nbsp;value while lower hydrogen ion activity corresponds to higher pH value. For pure water, its hydrogen ion concentration is &nbsp;10^(-7) which corresponds to pH of 7. So you can see that, pH is taken from the <em>power of 10</em>. For acidic solution, it has lower pH due to it having greater ion concentration, for example in the order of 10^(-4), which correspond to pH of &nbsp;4. pH can be better visualised in a scale as below:</p>\n<p><img src=\"https://i.imgsafe.org/ba/ba4a3ec56c.jpeg\" width=\"363\" height=\"169\"/></p>\n<p>That's all for now. I hope all of you can learn something from here. There are more theories on corrosion but i will post them in multiple different posts.</p>\n<p>Thank you.</p>\n</html>",
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